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ΔE with the PV work already handled. invented by someone who was tired. deeply relatable.
Finding ΔE requires measuring both q AND w every time. Most chemistry happens at constant atmospheric pressure. Measuring PV work every single experiment is annoying. Chemists, famously, did not want to do this.
So they defined enthalpy (H = E + PV) — a quantity where the PV term is silently absorbed into the definition. At constant pressure, ΔH = qp. You measure heat. You get enthalpy. You're done. The work problem simply no longer exists to you. Chemists looked at PV work and chose not to interact with it directly. Iconic.
| Process | ΔH | What's Happening | The Energy |
|---|---|---|---|
| Exothermic | ΔH < 0 | Heat exits system → warms surroundings. Flask heats up. | 🔥 combustion arc. things are on fire or functionally on fire. |
| Endothermic | ΔH > 0 | Heat enters system ← cools surroundings. Flask cools down. | 🧊 absorbing. pulling heat from the environment like a void with a plan. |
Double the moles → double the ΔH. The universe charges per mole and there is no bulk discount, no loyalty tier, no promotional pricing. CH₄ combustion is −890 kJ/mol. Two moles is −1780 kJ. It scales linearly and it is merciless about this.
Forward: ΔH = −890 kJ. Backward: ΔH = +890 kJ. Same number, opposite sign. This is non-negotiable. The universe has never once let someone reverse a reaction and keep the same ΔH. It will not start for you. Flip the equation, flip the sign. This is the law.
H₂O(l): ΔH°f = −285.8 kJ/mol. H₂O(g): ΔH°f = −241.8 kJ/mol. That is a 44 kJ/mol difference per mole of water. If you have 3 moles of water and write the wrong phase, you are off by 132 kJ and your answer is just wrong in a very specific and knowable way. Write the phase. (l) or (g). Every time.